Degrees of Freedom

1. Degrees of Freedom • Particles may have one or several types of freedom of motion. • and various degrees of each type • Translational freedom is the ability to move from one position in space to another. • Rotational freedom is the ability to reorient the particle’s direction in space. • Vibrational freedom is the ability to oscillate about a particular point in space.

2. The Structure of Gases • The particles in a gas are constantly in motion. • translational, vibrational, and rotational • There is a large amount of space between the particles. • compared to the size of the particles • Therefore, the molar volume of the gas state of a material is much larger than the molar volume of the solid or liquid states.

3. The Structure of Solids • The particles in a solid are packed tightly together. • The particles have no translational or rotational freedom of motion. • In crystalline solids, the particles are arranged in a geometric pattern. • In amorphous solids, the particles are stuck together without any long-range pattern.

4. The Structure of Liquids • The particles in a liquid are in close contact with each other (like a solid), but they have some limited ability to move around (like a gas). • The particles have limited translational and rotational freedom. • as well as complete vibrational freedom

5. Kinetic–Molecular Theory • The state of a material depends largely on two major factors: • the amount of kinetic energy the particles possess • the strength of attraction between the particles • These two factors are in competition with each other.

6. States and Degrees of Freedom • The molecules in a gas have complete freedom of motion. • Their kinetic energy overcomes the attractive forces between the molecules. • The molecules in a solid are locked in place; they cannot move around. • Though they do vibrate, they don’t have enough kinetic energy to overcome the attractive forces. • The molecules in a liquid have limited freedom—they can move around a little within the structure of the liquid. • They have enough kinetic energy to overcome some of the attractive forces, but not enough to escape each other.

7. Kinetic Energy • Increasing kinetic energy increases the motion energy of the particles. • The more motion energy the molecules have, the more freedom they can have. • The average kinetic energy is directly proportional to the temperature. • KEavg = 1.5 kT

8. Attractive Forces • The particles are attracted to each other by electrostatic attractive forces. • The strength of the attractive forces vary; some are small and some are large. • The strength of the attractive forces depends on the kind(s) of particles. • The stronger the attractive forces between the particles, the more they resist moving. • though no material completely lacks particle motion

9. Kinetic–Molecular Theory of Gases • When the attractive forces are much weaker than the kinetic energy, the material will be a gas. • In an ideal gas, the particles have complete freedom of motion—especially translational. • This allows gas particles to expand to fill their container. • Gases flow. • It also leads to large spaces between the particles. • therefore, low density and compressibility

10. Kinetic–Molecular Theory of Solids • When the attractive forces are strong enough so the kinetic energy cannot overcome it at all, the material will be a solid. • In a solid, the particles are packed together without any translational or rotational motion. • The only freedom they have is vibrational motion.

11. Explaining the Properties of Solids • The close packing of the particles results in solids being incompressible and having high density. • The inability of the particles to move around results in solids retaining their shape and volume when placed in a new container. • It also prevents the particles from flowing.

12. Solids • Some solids have their particles arranged in an orderly geometric pattern—we call these crystalline solids. • salt and diamonds • Other solids have particles that do not show a regular geometric pattern over a long range—we call these amorphous solids. • plastic and glass

13. Kinetic–Molecular Theoryof Liquids • When the attractive forces are strong enough so the kinetic energy can only partially overcome them, the material will be a liquid. • In a liquid, the particles are packed together with only very limited translational or rotational freedom.

14. Explaining the Properties of Liquids • Liquids have higher densities than gases and are incompressible because the particles are in contact. • They have an indefinite shape because the limited translational freedom of the particles allows them to move around enough to get to the container walls. • It also allows them to flow. • However, they have a definite volume because the limit on their freedom keeps them from escaping the rest of the particles.

15. Phase Changes • Since the attractive forces between the molecules are fixed, to change the material’s state requires changing the amount of kinetic energy the particles have, or limiting their freedom. • Solids melt when heated because the particles gain enough kinetic energy to partially overcome the attactive forces. • Liquids boil when heated because the particles gain enough kinetic energy to completely overcome the attractive forces. • The stronger the attractive forces, the higher you will need to raise the temperature. • Gases can be condensed by decreasing their temperature and/or increasing the pressure. • Pressure can be increased by decreasing the gas volume. • Reducing the volume reduces the amount of translational freedom the particles have.

16. Phase Changes

17. Why are molecules attracted to each other? • Intermolecular attractions are due to attractive forces between opposite charges. • + ion to − ion • + end of polar molecule to − end of polar molecule • H-bonding especially strong • Even nonpolar molecules will have temporary charges. • larger the charge = stronger attraction • longer the distance = weaker attraction • However, these attractive forces are small relative to the bonding forces between atoms. • generally smaller charges • generally over much larger distances

18. Trends in the Strength of Intermolecular Attraction • The stronger the attractions between the atoms or molecules, the more energy it will take to separate them. • Boiling a liquid requires addition of enough energy to overcome all the attractions between the particles. • however, not breaking the covalent bonds • The higher the normal boiling point of the liquid, the stronger the intermolecular attractive forces.

19. Kinds of Attractive Forces • Temporary polarity in the molecules due to unequal electron distribution leads to attractions called dispersion forces. • Permanent polarity in the molecules due to their structure leads to attractive forces called dipole–dipole attractions. • An especially strong dipole–dipole attraction results when H is attached to an extremely electronegative atom. These are called hydrogen bonds.

20. Dispersion Forces • Fluctuations in the electron distribution in atoms and molecules result in a temporary dipole. • A region with excess electron density has partial (–) charge. • A region with depleted electron density has partial (+) charge. • The attractive forces caused by these temporary dipoles are called dispersion forces. • aka London forces • All molecules and atoms will have them. • As a temporary dipole is established in one molecule, it induces a dipole in all the surrounding molecules.

22. + + + + + + + + + + + + + + + + - - - + + + + - - + + + + + + - - - - − − − − - − − − − − − − - - − − Size of the Induced Dipole • The magnitude of the induced dipole depends on several factors. • polarizability of the electrons • volume of the electron cloud • larger molar mass = more electrons = larger electron cloud = increased polarizability = stronger attractions • shape of the molecule • more surface-to-surface contact = larger induced dipole = stronger attraction Molecules that are flat have more surface interaction than spherical ones. Larger molecules have more electrons, leading to increased polarizability.

23. Dipole–Dipole Attractions • Polar molecules have a permanent dipole. • because of bond polarity and shape • dipole moment • as well as the always present induced dipole • The permanent dipole adds to the attractive forces between the molecules. • raising the boiling and melting points relative to nonpolar molecules of similar size and shape

24. Attractive Forces and Solubility • Solubility depends, in part, on the attractive forces of the solute and solvent molecules. • like dissolves like • Miscible liquids will always dissolve in each other. • Polar substances dissolve in polar solvents. • hydrophilic groups= OH, CHO, C=O, COOH, NH2, Cl • Nonpolar molecules dissolve in nonpolar solvents. • hydrophobic groups = C–H, C–C • Many molecules have both hydrophilic and hydrophobic parts. Solubility in water becomes a competition between the attraction of the polar groups for the water and the attraction of the nonpolar groups for their own kind.

25. Immiscible Liquids

26. Dichloromethane (methylene chloride) Ethanol (ethyl alcohol) Water Polar Solvents

27. Nonpolar Solvents n-pentane toluene carbon tetrachloride

28. Hydrogen Bonding • When a very electronegative atom is bonded to hydrogen, it strongly pulls the bonding electrons toward it. • O─H, N─H, or F─H • Since hydrogen has no other electrons, when it loses the electrons, the nucleus becomes deshielded. • exposing the H proton • The exposed proton acts as a very strong center of positive charge, attracting all the electron clouds from neighboring molecules.

29. H-Bonding HF

30. H-Bonding in Water

31. For nonpolar molecules, like the hydrides of group 4, the intermolecular attractions are due to dispersion forces. Therefore, they increase down the column causing the boiling point to increase. HF, H2O, and NH3 have unusally strong dipole–dipole attractions called hydrogen bonds. Therefore, they have higher boiling points than you would expect from the general trends. Polar molecules, like the hydrides of groups 5–7, have both dispersion forces and dipole–dipole attractions. Therefore, they have higher boiling points than the corresponding group 4 molecules.

32. Ion–Dipole Attraction • In a mixture, ions from an ionic compound are attracted to the dipole of polar molecules. • The strength of the ion–dipole attraction is one of the main factors that determines the solubility of ionic compounds in water.

33. Summary • Dispersion forces are the weakest of the intermolecular attractions. • Dispersion forces are present in all molecules and atoms. • The magnitude of the dispersion forces increases with molar mass. • Polar molecules also have dipole–dipole attractive forces.

34. Summary (cont’d) • Hydrogen bonds are the strongest of the intermolecular attractive forces. • a pure substance can have • Hydrogen bonds will be present when a molecule has H directly bonded to either O , N, or F atoms. • only example of H bonded to F is HF • Ion–dipole attractions are present in mixtures of ionic compounds with polar molecules. • Ion–dipole attractions are the strongest intermolecular attraction. • Ion–dipole attractions are especially important in aqueous solutions of ionic compounds.

35. Liquids Properties and Structure

36. Surface Tension • Surface tension is a property of liquids that results from the tendency of liquids to minimize their surface area. • In order to minimize their surface area, liquids form drops that are spherical. • as long as there is no gravity

37. Surface Tension • The layer of molecules on the surface behaves differently than the interior. • because the cohesive forces on the surface molecules have a net pull into the liquid interior • The surface layer acts like an elastic skin. • allowing you to “float” a fishing lure even though steel is denser than water

38. Surface Tension • Because they have fewer neighbors to attract them, the surface molecules are less stable than those in the interior. • have a higher potential energy • The surface tension of a liquid is the energy required to increase the surface area a given amount. • surface tension of H2O = 72.8 mJ/m2 • at room temperature • surface tension of C6H6 = 28 mJ/m2

39. Factors Affecting Surface Tension • The stronger the intermolecular attractive forces, the higher the surface tension will be. • Raising the temperature of a liquid reduces its surface tension. • Raising the temperature of the liquid increases the average kinetic energy of the molecules. • The increased molecular motion makes it easier to stretch the surface.

40. Viscosity • Viscosity is the resistance of a liquid to flow. • 1 poise = 1 P = 1 g/cm∙s • often given in centipoise, cP • H2O = 1 cP at room temperature • larger intermolecular attractions = larger viscosity

41. Factors Affecting Viscosity • The stronger the intermolecular attractive forces, the higher the liquid’s viscosity will be. • The more spherical the molecular shape, the lower the viscosity will be. • Molecules roll more easily. • Less surface-to-surface contact lowers attractions. • Raising the temperature of a liquid reduces its viscosity. • Raising the temperature of the liquid increases the average kinetic energy of the molecules. • The increased molecular motion makes it easier to overcome the intermolecular attractions and flow.

42. Capillary Action • Capillary action is the ability of a liquid to flow up a thin tube against the influence of gravity. • The narrower the tube, the higher the liquid rises. • Capillary action is the result of the two forces working in conjunction—the cohesive and adhesive forces. • Cohesive forces hold the liquid molecules together. • Adhesive forces attract the outer liquid molecules to the tube’s surface.

43. Capillary Action • The adhesive forces pull the surface liquid up the side of the tube, while the cohesive forces pull the interior liquid with it. • The liquid rises up the tube until the force of gravity counteracts the capillary action forces. • The narrower the tube diameter, the higher the liquid will rise up the tube.

44. Meniscus • The curving of the liquid surface in a thin tube is due to the competition between adhesive and cohesive forces. • The meniscus of water is concave in a glass tube because its adhesion to the glass is stronger than its cohesion for itself. • The meniscus of mercury is convex in a glass tube because its cohesion for itself is stronger than its adhesion for the glass. • Metallic bonds are stronger than intermolecular attractions.

45. The Molecular Dance • Molecules in the liquid are constantly in motion. • vibrational, and limited rotational and translational • The average kinetic energy is proportional to the temperature. • However, some molecules have more kinetic energy than the average, and others have less.

46. Vaporization • If these high-energy molecules are at the surface, they may have enough energy to overcome the attractive forces. • Therefore, the larger the surface area, the faster the rate of evaporation. • This will allow them to escape the liquid and become a vapor.

47. Distribution of Thermal Energy • Only a small fraction of the molecules in a liquid have enough energy to escape. • However, as the temperature increases, the fraction of the molecules with “escape energy” increases. • The higher the temperature, the faster the rate of evaporation.

48. Condensation • Some molecules of the vapor will lose energy through molecular collisions. • The result will be that some of the molecules will get captured back into the liquid when they collide with it. • Also, some may stick and gather together to form droplets of liquid. • particularly on surrounding surfaces • We call this process condensation.

49. Evaporation vs. Condensation • Vaporization and condensation are opposite processes. • In an open container, the vapor molecules generally spread out faster than they can condense. • The net result is that the rate of vaporization is greater than the rate of condensation, and there is a net loss of liquid. • However, in a closed container, the vapor is not allowed to spread out indefinitely. • The net result in a closed container is that at some time the rates of vaporization and condensation will be equal.

50. Effect of Intermolecular Attraction on Evaporation and Condensation • The weaker the attractive forces between molecules, the less energy they will need to vaporize. • Also, weaker attractive forces means that more energy will need to be removed from the vapor molecules before they can condense. • The net result will be more molecules in the vapor phase, and a liquid that evaporates faster—the weaker the attractive forces, the faster the rate of evaporation. • Liquids that evaporate easily are said to be volatile. • e.g., gasoline, fingernail polish remover • Liquids that do not evaporate easily are called nonvolatile. • e.g., motor oil